Atomic
Structure
Chapter 4
Introduction
• The view that we have of the atom has
been the result of work done over many
years by many people.
• The structure of the atom is composed of
smaller subatomic particles.
• Atoms can be distinguished from one
another based upon their atomic and mass
numbers.
• Isotopes are different versions of a given
atom and contribute to the average mass
recorded in the periodic table.
Defining the Atom
(Section 4.1)
• Early
Models of
the Atom
• Sizing Up
The Atom
The Macro vs. The Micro
What are some differences?
What are some similarities?
Elements
• Elements: The simplest form of matter
that has a unique set of properties (i.e.
different physical and chemical
properties). [Section 2.3]
In other words all elements behave
differently.
Elements = Atoms
These two terms are
used interchangeably.
Therefore, when we talk about an atom
of oxygen, we are referring to the element
oxygen.
The Atom
The smallest particle of an element
that retains its identity in a
chemical reaction.
The word comes from the Greek word atomos
which means “something that cannot be divided
further.”
The radii of most atoms fall within the range of
5 x 10-11 m to 2 x 10-10m (50 pm – 200 pm).
STM image of gold atoms
STM image of iron atoms
STM image of platinum atoms
I.) Early Models of the Atom
• Democritus at about 400
B.C. proposed that all
things were composed of
extremely small and
irreducible particles.
– He lacked experimental
support for his idea and
his approach was not
based on the scientific
method.
These ideas were rejected by Aristotle and were
consequently forgotten for almost 2000 years.
• 1803 – John Dalton,
an English chemist
and school teacher,
proposed the modern
version of the atomic
theory.
– Dalton used
experimental methods
and transformed
Democritus’s ideas on
atoms into a scientific
theory
Dalton studied the ratios in which elements
combine in chemical reactions.
Dalton’s Atomic Theory
1. All elements are composed of tiny
indivisible particles called atoms.
Oxygen atom
Hydrogen atoms
2. Atoms of the same element are identical.
The atoms of any one element are
different from those of any other
element.
STM Image of Fe
STM image of Au
3. Atoms of different elements can
physically mix together or can
chemically combine in simple wholenumber ratios to form compounds.
4.Chemical reactions occur when atoms are
separated, joined, or rearranged. Atoms of
one element are never changed into atoms
of another element as a result of a
chemical reaction.
The electolysis of water
II.) Sizing Up The Atom
Scanning Tunneling Microscope (STM)
How a STM works.
Xenon on Nickel
Carbon Monoxide
on Platinum
Is There Anything
Smaller Than the Atom?
Yes.
These are called subatomic particles.
Structure of the Nuclear Atom
(Section 4.2)
• Subatomic
Particles
• The Atomic
Nucleus
I.) Subatomic Particles
The idea of a nuclear atom showed that
atoms are divisible.
Subatomic Particles
Electrons:
negatively charged
subatomic particles.
Protons: positively
charged
subatomic particles.
Neutrons:
subatomic particles
with no charge.
Characteristics of the Subatomic
Particles
Discovering the Electron
• In 1897 uncovered the existence of
electrons using cathode ray tubes
(CRT).
• Hypothesized that a cathode ray is a
stream of negatively charged particles.
• He set up an experiment to measure
the ratio of the charge of an electron to
its mass. This ratio does not depend
on the kind of gas in the CRT.
J.J. Thomson
• Robert Millikan, an American physicist English Physicist
(1856 - 1940)
used this ratio to calculate the mass of
the electron.
Thomson's Experiment
CRT with no magnet
deflecting
stream
CRT with magnet
deflecting stream
If electrons are uncharged,
then no deflection would
occur at all.
II.) The Atomic Nucleus
• Thomson proposed the "plum pudding" model of
the atom in which the electrons are stuck into a
lump of positive charge.
• No nucleus in this model.
Discovering the Proton
• In 1886, Goldstein found
rays travelling in the
direction opposite of the
cathode rays. He called
these rays "canal rays."
• He believed these rays
were composed of
positive particles. We
now know these
particles to be protons.
Eugene Goldstein
German Physicist
(1850 - 1930)
The Rutherford Atom
• A former student of J.J
Thomson.
• Is considered the "father" of
nuclear physics.
• Devised the planetary model
of the atom, which is Neils
Bohr's orbital model.
• Through experiments
looking at the scattering
pattern of certain particles,
he reached his conclusions.
Ernest Rutherford
(1871 - 1937)
Rutherford's Gold-Foil
Experiment
• Experiment used alpha particles, which
are helium atoms with their electrons
removed, thus they have two positive
charges.
• The alpha particles are shot at a thin sheet
of gold foil (the cannon ball shot at a tissue
paper).
• The alpha particles then strike a
fluorescent screen behind the gold foil.
Interpreting Rutherford's Results
• If the plum pudding model was true, then almost all
the alpha particles would have passed through the
foil with only some slight deflection.
• The fact that some particles bounced back indicate
that they collided with a similarly charged particle.
(The tissue paper sending the cannon ball back at
the cannon.
• Conclsion:
1. Atom mostly space
2. At the center some positive charge exists that
accounts for the deflection.
3. This positively charged region he termed the
nucleus.
• The particles did
not all pass
through.
• The deflection of
the positive
alpha particles
shows there is a
positive charge
in the atom.
The Rutherford Model of the
Atom
• A centrally located nucleus that
contains almost all the mass of
the atom and all the positive
charge.
• Electrons are negatively
charged and distributed around
the nucleus. The electrons'
orbits accounts for almost all
the volume of the atom.
• This model is incomplete and
fails to explain the chemical
behavior of elements. It needed
to be improved upon.
Discovering the Neutron
• Worked in Ernest Rutherford's
lab before and after WWI.
• Discovered the neutron while
studying atomic disintegration.
He was looking for a particle
with a mass equal to the
proton but possessed no
charge.
• Werner Heisenberg, using
Chadwick's discovery, showed
that neutrons were not simply
electron/proton pairs, but
their own particles.
James Chadwick
English Physicist
(1891 - 1974)
Label the diagram of the atom
Distinguishing Among Atoms
(Section 4.3)
•
•
•
•
•
Atomic Number
Mass Number
Isotopes
Atomic Mass
The Periodic
Table – A
Preview
I.) Atomic Number & Mass
Number
• Elements are different from one
another because they contain
different numbers of protons.
• Atomic Number: The number of
protons in the nucleus of an atom.
• Mass Number: The total number of
neutrons and protons in an atom.
Each of these numbers can be found
for each element using the periodic
table.
The mass number is always the bigger number.
Calculating the number of
neutrons in an atom.
# of Neutrons = Mass number atomic number
• Using the atomic number, the
mass number, and the above
equation we can determine the
number of protons, neutrons, and
electrons for any atom.
Representing Atomic and
Mass Numbers
• Use the symbol for the
element
• Place the mass number
as a superscript on the
left-hand side of the
symbol.
• Place the atomic
number as the subscript
on the left-hand side of
the symbol
197
79
Au
Gold-197
II.) Isotopes
Atoms that have the same number of
protons but different numbers of
neutrons.
• Isotopes are chemically similar, but
not identical.
• Isotopes of a given element have
different mass numbers.
• Each element has a specific number
of isotopes (i.e. hydrogen has three
naturally occurring isotopes).
• Not all isotopes occur with the same
abundance in nature.
Determining the Composition of
an Isotope
• The composition of an isotope is
determined in the same way as an
element.
3 isotopes of chromium exists:
chromium-50, chromium-52, and
chromium-53. How many
neutrons, protons, and electrons
are in each isotope, given that the
atomic number is 24?
III.) Atomic Masses
• Actual atomic masses are very small
– F = 3.155 x 10-23 g.
– As = 1.244 x 10-22 g.
• Numbers this small are impractical to work
with.
• Better to use relative masses of atoms
using carbon-12 as a reference standard.
– Carbon-12 = 12 amu.
Atomic Mass Units (amu)
• 1 amu (u) is defined as one-twelfth the
mass of a carbon -12 atom.
• 1 amu = 1.66 x 10-24 g.
• Helium = 4.0026 u
Carbon = 12.011 u
If the masses of atoms depend primarily
on the number of protons and neutrons,
why aren’t the masses listed in the
periodic table whole numbers?
In nature most elements occur as a mixture of
two or more isotopes, each having a different
mass.
• Remember, isotopes of an element have
different numbers of neutrons and
thus will have different masses.
Natural Percent Abundance of Stable Isotopes
of Some Elements
Name
Hydrogen
Symbol
1H
2H
3H
Carbon
12C
13C
Oxygen
16O
17O
18O
Natural %
Abundance
Mass
(amu)
99.985 %
0.015 %
negligible
1.0078 u
2.0141 u
3.0160 u
98.89 %
1.11 %
12.000 u
13.003 u
99.759 %
0.037 %
0.204 %
15.995 u
16.995 u
17.999 u
Average
atomic
mass
1.0079
12.011
15.999
Atomic Masses
• The masses listed in the periodic table are
weighted average of the masses of all the
elements.
• This weighted average is known as the atomic
mass.
– A weighted average mass of the atoms in a
naturally occurring sample of an element.
• A weighted mass reflects both the mass and the
relative abundance of the isotopes of an
element.
Calculating Atomic Masses
•
To do this calculation you need to know:
– the number of stable isotopes
– the mass of each isotope (mi)
– the natural percentage abundance of each
isotope (%)
Atomic mass = Σ[(mi) x %]
Let’s practice
Determine the atomic mass of carbon.
Carbon has two stable isotopes.
Carbon-12 = 12.000 u (98.99%)
Carbon-13 = 13.003 u (1.11%)
IV.) The Periodic Table – A
Preview
• The periodic table is an arrangement of
elements in which the elements are
separated into groups based on a set of
repeating properties.
• A periodic table allows us to easily
compare the properties of one element (or
group of elements) to another element (or
group of elements).
We will discuss the periodic table
in greater detail in chapter 6.
(7 periods)
(18 periods)
The periodic table gives a wealth of info.
regarding the elements. Such as:
• Atomic and mass numbers.
• Chemical properties of each element.
• The sizes of the atoms.
• The location of electrons in each atom.
Use the periodic table to locate the following
elements, given their atomic numbers.
1. Atomic Number = 29
2. Atomic Number = 50
3. Atomic Number = 82
4. Atomic Number = 79
Now write their chemical symbols. Include the
atomic number and mass number.
Atomic
Structure
Chapter 4
The End